Lewis Structures:
if it’s a halogen, with 7 valence electrons, it can only form one bond
if it has 6 valence electrons, it can form a double bond
if it has 5 valence electrons it can form a triple bond
if it has 4 valence electrons it can form a triple bond as well
the atom has to be responsible for as many indivisual electrons as there are valence electrons
sulfur and phosphorus are exceptions because they can surpass the octet rule; this is because they have empty d orbitals that can be filled with extra electrons, giving rise to the possibility of an expanded octet
Formal charge:
Valence electrons-electrons in lone pairs-# of bonds
the goal is for the formal charge to be zero on most of the atoms
each charge adds to the overall charge
Resonance structures:
molecular compounds with more than 1 possible lewis structure
VSEPR Shapes
linear if there are 2 atoms bonded to the central atom and zero lone pairs
180 degrees
linear if there are only 2 atoms; can’t make a shape
180 degrees
bent if there are 2 atoms bonded to the central atom and one lone pair
a little <109.5 degrees
bent if there are 2 atoms bonded to the central atom and 2 lone pairs
a little <109.5 degrees
trigonal planar if there are 3 atoms bonded to the central atom and 0 lone pairs
120 degrees
trigonal pyramidal if there are 3 atoms bonded to the central atom and 1 lone pair
a little <109.5 degrees
tetrahedral if there are 4 atoms bonded to the central atom and 0 lone pairs
109.5 degrees
Electronegativity and Intermolecular Forces
types of bonds
london dispersion, in all molecules
dipole-dipole, in polar molecules
hydrogen bonding in polar molecules with H-F, H-O, H-N bonds
compound is nonpolar if
the central atom has 0 lone pairs
all atoms around central atom are the same
dipole moments point toward the most electronegative molecules
bond polarity is toward the more electronative atom
molecular polarity is the overall dipole moment
pauling scale
If the electronegativity is
<0.4, nonpolar
0.4-1.7, polar
>1.7, ionic